CHAPTER. Copyright by Holt, Rinehart and Winston. All rights reserved.

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1 CHAPTER 72

2 Until recently, if you wanted to see an image of atoms, the best you could hope to see was an artists s drawing of atoms. Now, with the help of powerful microscopes, scientists are able to obtain images of atoms. One such microscope is known as the scanning tunneling microscope, which took the image of the nickel atoms shown on the opposite page. As its name implies, this microscope scans a surface, and it can come as close as a billionth of a meter to a surface to get an image. The images that these microscopes provide help scientists understand atoms. START-UPACTIVITY Forces of Attraction PROCEDURE SAFETY PRECAUTIONS 1. Spread some salt and pepper on a piece of paper that lies on a flat surface. Mix the salt and pepper but make sure that the salt and pepper are not clumped together. 2. Rub a plastic spoon with a wool cloth. 3. Hold the spoon just above the salt and pepper. 4. Clean off the spoon by using a towel. Rub the spoon with the wool cloth and bring the spoon slowly toward the salt and pepper from a distance. ANALYSIS 1. What happened when you held your spoon right above the salt and pepper? What happened when you brought your spoon slowly toward the salt and pepper? 2. Why did the salt and pepper jump up to the spoon? 3. When the spoon is brought toward the paper from a distance, which is the first substance to jump to the spoon? Why? CONTENTS SECTION 1 Substances Are Made of Atoms SECTION 2 Structure of Atoms SECTION 3 Electron Configuration SECTION 4 Counting Atoms Pre-Reading Questions What is an atom? What particles make up an atom? Where are the particles that make up an atom located? Name two types of electromagnetic radiation. Topic: Atoms and Elements SciLinks code: HW

3 SECTION 1 Substances Are Made of Atoms KEY TERMS law of definite proportions law of conservation of mass law of multiple proportions Topic: Development of Atomic Theory SciLinks code: HW OBJECTIVES State the three laws that support the existence of atoms. List the five principles of John Dalton s atomic theory. Atomic Theory As early as 400 BCE, a few people believed in an atomic theory, which states that atoms are the building blocks of all matter. Yet until recently, even scientists had never seen evidence of atoms. Experimental results supporting the existence of atoms did not appear until more than 2000 years after the first ideas about atoms emerged. The first of these experimental results indicated that all chemical compounds share certain characteristics. What do you think an atom looks like? Many people think that an atom looks like the diagram in Figure 1a. However, after reading this chapter, you will find that the diagram in Figure 1b is a better model of an atom. Recall that a compound is a pure substance composed of atoms of two or more elements that are chemically combined. These observations about compounds and the way that compounds react led to the development of the law of definite proportions, the law of conservation of mass, and the law of multiple proportions. Experimental observations show that these laws also support the current atomic theory. Topic : Current Atomic Theory SciLinks code: HW4038 Figure 1 a Many people believe that an atom looks like this diagram. b This diagram is a better model of the atom. 74 Chapter 3

4 The Law of Definite Proportions The law of definite proportions states that two samples of a given compound are made of the same elements in exactly the same proportions by mass regardless of the sizes or sources of the samples. Notice the composition of ethylene glycol, as shown in Figure 2. Every sample of ethylene glycol is composed of three elements in the following proportions by mass: law of definite proportions the law that states that a chemical compound always contains the same elements in exactly the same proportions by weight or mass 51.56% oxygen, 38.70% carbon, and 9.74% hydrogen The law of definite proportions also states that every molecule of ethylene glycol is made of the same number and types of atoms. A molecule of ethylene glycol has the formula C 2 H 6 O 2, so the law of definite proportions tells you that all other molecules of ethylene glycol have the same formula. Table salt (sodium chloride) is another example that shows the law of definite proportions. Any sample of table salt consists of two elements in the following proportions by mass: 60.66% chlorine and 39.34% sodium Every sample of table salt also has the same proportions of ions. As a result, every sample of table salt has the same formula, NaCl. As chemists of the 18th century began to gather data during their studies of matter, they first began to recognize the law of definite proportions. Their conclusions led to changes in the atomic theory. STUDY TIP USING THE ILLUSTRATIONS The illustrations in the text will help you make the connection between what you can see, such as a beaker of chemicals, and what you cannot see, such as the atoms that make up those chemicals. Notice that the model in Figure 2 shows how the atoms of a molecule of ethylene glycol are arranged. To practice thinking at the particle level, draw pictures of water molecules and copper atoms. Figure 2 a Ethylene glycol is the main component of automotive antifreeze. b Ethylene glycol is composed of carbon, oxygen, and hydrogen. c Ethylene glycol is made of exact proportions of these elements regardless of the size of the sample or its source. Ethylene Glycol Composition by Mass oxygen 51.56% carbon 38.70% hydrogen 9.74% Atoms and Moles 75

5 law of conservation of mass the law that states that mass cannot be created or destroyed in ordinary chemical and physical changes The Law of Conservation of Mass As early chemists studied more chemical reactions, they noticed another pattern. Careful measurements showed that the mass of a reacting system does not change. The law of conservation of mass states that the mass of the reactants in a reaction equals the mass of the products. Figure 3 shows several reactions that show the law of conservation of mass. For example, notice the combined mass of the sulfur atom and the oxygen molecule equals the mass of the sulfur dioxide molecule. Also notice that Figure 3 shows that the sum of the mass of the chlorine molecule and the mass of the phosphorus trichloride molecule is slightly smaller than the mass of the phosphorus pentachloride molecule. This difference is the result of rounding off and of correctly using significant figures. Figure 3 The total mass of a system remains the same whether atoms are combined, separated, or rearranged. Here, mass is expressed in kilograms (kg). Conservation of Mass Hydrogen molecule kg Oxygen atom kg Water molecule kg 1 2 H 2 + O 2 H 2 O Sulfur atom kg Oxygen molecule kg Sulfur dioxide molecule kg S + O 2 SO 2 Phosphorus pentachloride molecule kg Phosphorus trichloride molecule kg Chlorine molecule kg PCl 5 PCl 3 + Cl 2 76 Chapter 3

6 Table 1 Compounds of Nitrogen and Oxygen and the Law of Multiple Proportions Name of Description As shown Formula Mass Mass compound in figures O (g) N (g) M ass O( g) Mass N( g) Nitrogen colorless gas monoxide that reacts readily NO g O = g O with oxygen g N 1 g N Nitrogen poisonous dioxide brown gas NO g O = g O in smog g N 1 g N The Law of Multiple Proportions Table 1 lists information about the compounds nitrogen monoxide and nitrogen dioxide. For each compound, the table also lists the ratio of the mass of oxygen to the mass of nitrogen. So, 1.14 g of oxygen combine with 1 g of nitrogen when nitrogen monoxide forms. In addition, 2.28 g of oxygen combine with 1 g of nitrogen when nitrogen dioxide forms. The ratio of the masses of oxygen in these two compounds is exactly 1.14 to 2.28 or 1 to 2. This example illustrates the law of multiple proportions: If two or more different compounds are composed of the same two elements, the ratio of the masses of the second element (which combines with a given mass of the first element) is always a ratio of small whole numbers. The law of multiple proportions may seem like an obvious conclusion given the molecules diagrams and formulas shown. But remember that the early chemists did not know the formulas for compounds. In fact, chemists still have not actually seen these molecules. Scientists think that molecules have these formulas because of these mass data. law of multiple proportions the law that states that when two elements combine to form two or more compounds, the mass of one element that combines with a given mass of the other is in the ratio of small whole numbers Dalton s Atomic Theory In 1808, John Dalton, an English school teacher, used the Greek concept of the atom and the law of definite proportions, the law of conservation of mass, and the law of multiple proportions to develop an atomic theory. Dalton believed that a few kinds of atoms made up all matter. According to Dalton, elements are composed of only one kind of atom and compounds are made from two or more kinds of atoms. For example, the element copper consists of only one kind of atom, as shown in Figure 4. Notice that the compound iodine monochloride consists of two kinds of atoms joined together. Dalton also reasoned that only whole numbers of atoms could combine to form compounds, such as iodine monochloride. In this way, Dalton revised the early Greek idea of atoms into a scientific theory that could be tested by experiments. iodine monochloride copper Figure 4 An element, such as copper, is made of only one kind of atom. In contrast, a compound, such as iodine monochloride, can be made of two or more kinds of atoms. Atoms and Moles 77

7 Dalton s Theory Contains Five Principles Dalton s atomic theory can be summarized by the following statements: 1. All matter is composed of extremely small particles called atoms, which cannot be subdivided, created, or destroyed. 2. Atoms of a given element are identical in their physical and chemical properties. 3. Atoms of different elements differ in their physical and chemical properties. 4. Atoms of different elements combine in simple, whole-number ratios to form compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged but never created, destroyed, or changed. Dalton s theory explained most of the chemical data that existed during his time. As you will learn later in this chapter, data gathered since Dalton s time shows that the first two principles are not true in all cases. Today, scientists can divide an atom into even smaller particles. Technology has also enabled scientists to destroy and create atoms. Another feature of atoms that Dalton could not detect is that many atoms will combine with like atoms. Oxygen, for example, is generally found as O 2, a molecule made of two oxygen atoms. Sulfur is found as S 8. Because some parts of Dalton s theory have been shown to be incorrect, his theory has been modified and expanded as scientists learn more about atoms. 1 Section Review UNDERSTANDING KEY IDEAS 1. What is the atomic theory? 2. What is a compound? 3. State the laws of definite proportions, conservation of mass, and multiple proportions. 4. According to Dalton, what is the difference between an element and a compound? 5. What are the five principles of Dalton s atomic theory? 6. Which of Dalton s five principles still apply to the structure of an atom? CRITICAL THINKING 7. What law is described by the fact that carbon dioxide consists of 27.3% carbon and 72.7% oxygen by mass? 8. What law is described by the fact that the ratio of the mass of oxygen in carbon dioxide to the mass of oxygen in carbon monoxide is 2:1? 9. Three compounds contain the elements sulfur, S, and fluorine, F. How do the following data support the law of multiple proportions? compound A: g F for every g S compound B: g F for every g S compound C: g F for every g S 78 Chapter 3

8 SECTION 2 Structure of Atoms KEY TERMS electron nucleus proton neutron atomic number mass number isotope OBJECTIVES Describe the evidence for the existence of electrons, protons, and neutrons, and describe the properties of these subatomic particles. Discuss atoms of different elements in terms of their numbers of electrons, protons, and neutrons, and define the terms atomic number and mass number. Define isotope, and determine the number of particles in the nucleus of an isotope. Subatomic Particles Experiments by several scientists in the mid-1800s led to the first change to Dalton s atomic theory. Scientists discovered that atoms can be broken into pieces after all. These smaller parts that make up atoms are called subatomic particles. Many types of subatomic particles have since been discovered. The three particles that are most important for chemistry are the electron, the proton, and the neutron. Topic : Subatomic Particles SciLinks code: HW4121 Electrons Were Discovered by Using Cathode Rays The first evidence that atoms had smaller parts was found by researchers who were studying electricity, not atomic structure. One of these scientists was the English physicist J. J. Thomson. To study current, Thomson pumped most of the air out of a glass tube. He then applied a voltage to two metal plates, called electrodes, which were placed at either end of the tube. One electrode, called the anode, was attached to the positive terminal of the voltage source, so it had a positive charge. The other electrode, called a cathode, had a negative charge because it was attached to the negative terminal of the voltage source. Thomson observed a glowing beam that came out of the cathode and struck the anode and the nearby glass walls of the tube. So, he called these rays cathode rays. The glass tube Thomson used is known as a cathode-ray tube (CRT). CRTs have become an important part of everyday life. They are used in television sets, computer monitors, and radar displays. An Electron Has a Negative Charge Thomson knew the rays must have come from the atoms of the cathode because most of the atoms in the air had been pumped out of the tube. Because the cathode ray came from the negatively charged cathode, Thomson reasoned that the ray was negatively charged. Figure 5 The image on a television screen or a computer monitor is produced when cathode rays strike the special coating on the inside of the screen. Atoms and Moles 79

9 Figure 6 A magnet near the cathode-ray tube causes the beam to be deflected. The deflection indicates that the particles in the beam have a negative charge. magnet anode cathode deflected beam vacuum pump Topic : J. J. Thomson SciLinks code: HW4156 electron a subatomic particle that has a negative electric charge He confirmed this prediction by seeing how electric and magnetic fields affected the cathode ray. Figure 6 shows what Thomson saw when he placed a magnet near the tube. Notice that the beam is deflected by the magnet. Other researchers had shown that moving negative charges are deflected this way. Thomson also observed that when a small paddle wheel was placed in the path of the rays, the wheel would turn. This observation suggested that the cathode rays consisted of tiny particles that were hitting the paddles of the wheel. Thomson s experiments showed that a cathode ray consists of particles that have mass and a negative charge. These particles are called electrons. Table 2 lists the properties of an electron. Later experiments, which used different metals for cathodes, confirmed that electrons are a part of atoms of all elements. Electrons are negatively charged, but atoms have no charge. Therefore, atoms must contain some positive charges that balance the negative charges of the electrons. Scientists realized that positive charges must exist in atoms and began to look for more subatomic particles. Scientists also recognized that atoms must have other particles because an electron was found to have much less mass than an atom does. Table 2 Properties of an Electron Name Symbol As shown Charge Common charge Mass (kg) in figures notation Electron e, e,or 1 0 e C kg 80 Chapter 3

10 Rutherford Discovered the Nucleus Thomson proposed that the electrons of an atom were embedded in a positively charged ball of matter. His picture of an atom, which is shown in Figure 7, was named the plum-pudding model because it resembled plum pudding, a dessert consisting of a ball of cake with pieces of fruit in it. Ernest Rutherford, one of Thomson s former students, performed experiments in 1909 that disproved the plum-pudding model of the atom. Rutherford s team of researchers carried out the experiment shown in Figure 8. A beam of small, positively charged particles, called alpha particles, was directed at a thin gold foil.the team measured the angles at which the particles were deflected from their former straight-line paths as they came out of the foil. Rutherford found that most of the alpha particles shot at the foil passed straight through the foil. But a very small number of particles were deflected, in some cases backward, as shown in Figure 8. This result greatly surprised the researchers it was very different from what Thomson s model predicted.as Rutherford said, It was almost as if you fired a 15-inch shell into a piece of tissue paper and it came back and hit you. After thinking about the startling result for two years, Rutherford finally came up with an explanation. He went on to reason that only a very concentrated positive charge in a tiny space within the gold atom could possibly repel the fast-moving, positively charged alpha particles enough to reverse the alpha particles direction of travel. Rutherford also hypothesized that the mass of this positive-charge containing region, called the nucleus, must be larger than the mass of the alpha particle. If not, the incoming particle would have knocked the positive charge out of the way. The reason that most of the alpha particles were undeflected, Rutherford argued, was that most parts of the atoms in the gold foil were empty space. This part of the model of the atom is still considered true today. The nucleus is the dense, central portion of the atom. The nucleus has all of the positive charge, nearly all of the mass, but only a very small fraction of the volume of the atom. Figure 7 Thomson s model of an atom had negatively charged electrons embedded in a ball of positive charge. nucleus an atom s central region, which is made up of protons and neutrons Figure 8 Greatly deflected particle Beam of positively charged subatomic particles Slightly deflected particle Undeflected particles a In the gold foil experiment, small positively charged particles were directed at a thin foil of gold atoms. Nucleus of gold atom b The pattern of deflected alpha particles supported Rutherford s hypothesis that gold atoms were mostly empty space. Gold atom Atoms and Moles 81

11 Figure 9 If the nucleus of an atom were the size of a marble, then the whole atom would be about the size of a football stadium. proton a subatomic particle that has a positive charge and that is found in the nucleus of an atom; the number of protons of the nucleus is the atomic number, which determines the identity of an element neutron a subatomic particle that has no charge and that is found in the nucleus of an atom Protons and Neutrons Compose the Nucleus By measuring the numbers of alpha particles that were deflected and the angles of deflection, scientists calculated the radius of the nucleus to be 1 less than 10 of the radius of the whole atom. Figure 9 gives you a better 000 idea of these sizes. Even though the radius of an entire atom is more than times larger than the radius of its nucleus, an atom is still extremely small. The unit used to express atomic radius is the picometer (pm). One picometer equals m. The positively charged particles that repelled the alpha particles in the gold foil experiments and that compose the nucleus of an atom are called protons. The charge of a proton was calculated to be exactly equal in magnitude but opposite in sign to the charge of an electron. Later experiments showed that the proton s mass is almost 2000 times the mass of an electron. Because protons and electrons have equal but opposite charges, a neutral atom must contain equal numbers of protons and electrons. But solving this mystery led to another: the mass of an atom (except hydrogen atoms) is known to be greater than the combined masses of the atom s protons and electrons. What could account for the rest of the mass? Hoping to find an answer, scientists began to search for a third subatomic particle. About 30 years after the discovery of the electron, Irene Joliot-Curie (the daughter of the famous scientists Marie and Pierre Curie) discovered that when alpha particles hit a sample of beryllium, a beam that could go through almost anything was produced. The British scientist James Chadwick found that this beam was not deflected by electric or magnetic fields. He concluded that the particles carried no electric charge. Further investigation showed that these neutral particles, which were named neutrons, are part of all atomic nuclei (except the nuclei of most hydrogen atoms). 82 Chapter 3

12 Table 3 Properties of a Proton and a Neutron Name Symbol As shown Charge Common charge Mass (kg) in figures notation Proton p, p +, or p C kg Neutron n or 0 1 n 0 C kg Protons and Neutrons Can Form a Stable Nucleus Table 3 lists the properties of a neutron and a proton. Notice that the charge of a neutron is commonly assigned the value 0 while that of a proton is +1. How do protons that are positively charged come together to form a nucleus? In fact, the formation of a nucleus with protons seems impossible if you just consider Coulomb s law. Coulomb s law states that the closer two charges are, the greater the force between them. In fact, the force increases by a factor of 4 as the distance is halved. In addition, the larger the two charges are the greater the force between them. If the charges are opposite, they attract one another. If both charges have the same sign, they repel one another. If you keep Coulomb s law in mind, it is easy to understand why with the exception of some hydrogen atoms no atoms have nuclei that are composed of only protons.all protons have a +1 charge. So, the repulsive force between two protons is large when two protons are close together, such as within a nucleus. Protons, however, do form stable nuclei despite the repulsive force between them. A strong attractive force between these protons overcomes the repulsive force at small distances. Because neutrons also add attractive forces without being subject to repulsive charge-based forces, some neutrons can help stabilize a nucleus. Thus, all atoms that have more than one proton also have neutrons. Topic: Atomic Nucleus SciLinks Code: HW As charge increases, force of attraction increases Figure 10 This figure shows that the larger two charges are, the greater the force between the charges. In addition, the figure shows the smaller the distance between two charges, the greater the force between the charges. larger distance smaller distance As distance decreases, force of attraction increases Atoms and Moles 83

13 Atomic Number and Mass Number All atoms consist of protons and electrons. Most atoms also have neutrons. Protons and neutrons make up the small, dense nuclei of atoms. The electrons occupy the space surrounding the nucleus. For example, an oxygen atom has protons and neutrons surrounded by electrons. But that description fits all other atoms, such as atoms of carbon, nitrogen, silver, and gold. How, then, do the atoms of one element differ from those of another element? Elements differ from each other in the number of protons their atoms contain. atomic number the number of protons in the nucleus of an atom; the atomic number is the same for all atoms of an element Atomic Number Is the Number of Protons of the Nucleus The number of protons that an atom has is known as the atom s atomic number. For example, the atomic number of hydrogen is 1 because the nucleus of each hydrogen atom has one proton. The atomic number of oxygen is 8 because all oxygen atoms have eight protons. Because each element has a unique number of protons in its atoms, no two elements have the same atomic number. So an atom whose atomic number is 8 must be an oxygen atom. To date, scientists have identified 113 elements, whose atomic numbers range from 1 to 114.The element whose atomic number is 113 has yet to be discovered. Note that atomic numbers are always whole numbers. For example, an atom cannot have 2.5 protons. The atomic number also reveals the number of electrons in an atom of an element. For atoms to be neutral, the number of negatively charged electrons must equal the number of positively charged protons. Therefore, if you know the atomic number of an atom, you immediately know the number of protons and the number of electrons found in that atom. Figure 11 shows a model of an oxygen atom, whose atomic number is 8 and which has 8 electrons surrounding a nucleus that has 8 protons. The atomic number of gold is 79, so an atom of gold must have 79 electrons surrounding a nucleus of 79 protons. The next step in describing an atom s structure is to find out how many neutrons the atom has. Figure 11 The atomic number for oxygen, as shown on the periodic table, tells you that the oxygen atom has 8 protons and 8 electrons. Symbol of element Name of element Atomic number 8 O Oxygen [He]2s 2 2p 4 Proton Mass of element Electron configuration Electron cloud Neutron 84 Chapter 3

14 Mass Number Is the Number of Particles of the Nucleus Every atomic nucleus can be described not only by its atomic number but also by its mass number.the mass number is equal to the total number of particles of the nucleus that is, the total number of protons and neutrons. For example, a particular atom of neon has a mass number of 20, as shown in Figure 12. Therefore, the nucleus of this atom has a total of 20 protons and neutrons. Because the atomic number for an atom of neon is 10, neon has 10 protons. You can calculate the number of neutrons in a neon atom by subtracting neon s atomic number (the number of protons) from neon s mass number (the number of protons and neutrons). mass number the sum of the numbers of protons and neutrons of the nucleus of an atom mass number atomic number = number of neutrons In this example, the neon atom has 10 neutrons. number of protons and neutrons (mass number) = 20 number of protons (atomic number) = 10 number of neutrons = 10 Unlike the atomic number, which is the same for all atoms of an element, mass number can vary among atoms of a single element. In other words, all atoms of an element have the same number of protons, but they can have different numbers of neutrons. The atomic number of every hydrogen atom is 1, but hydrogen atoms can have mass numbers of 1, 2, or 3. These atoms differ from one another in having 0, 1, and 2 neutrons, respectively. Another example is oxygen. The atomic number of every oxygen atom is 8, but oxygen atoms can have mass numbers of 16, 17, or 18. These atoms differ from one another in having 8, 9, and 10 neutrons, respectively. Figure 12 The neon atom has 10 protons, 10 neutrons, and 10 electrons. This atom s mass number is 20, or the sum of the numbers of protons and neutrons in the atom. Proton Atomic number Symbol of element Name of element Electron configuration Mass of element Electron cloud Neutron Atoms and Moles 85

15 SAMPLE PROBLEM A PRACTICE HINT Check that the atomic number and the number of protons are the same. Also check that adding the numbers of protons and neutrons equals the mass number. Determining the Number of Particles in an Atom How many protons, electrons, and neutrons are present in an atom of copper whose atomic number is 29 and whose mass number is 64? 1 Gather information. The atomic number of copper is 29. The mass number of copper is Plan your work. The atomic number indicates the number of protons in the nucleus of a copper atom. A copper atom must be electrically neutral, so the number of electrons equals the number of protons. The mass number indicates the total number of protons and neutrons in the nucleus of a copper atom. 3 Calculate. atomic number (29) = number of protons = 29 number of protons = number of electrons = 29 mass number (64) atomic number (29) = number of neutrons = 35 4 Verify your results. number of protons (29) + number of neutrons (35) = mass number (64) PRACTICE PROBLEM SOLVING SKILL 1 How many protons and electrons are in an atom of sodium whose atomic number is 11? 2 An atom has 13 protons and 14 neutrons. What is its mass number? 3 Calculate the mass number for an atom that has 45 neutrons and 35 electrons. 4 An atom of an element has 54 protons. Some of the element s atoms have 77 neutrons, while other atoms have 79 neutrons. What are the atomic numbers and mass numbers of the two types of atoms of this element? Different Elements Can Have the Same Mass Number The atomic number identifies an element. For example, copper has the atomic number 29. All copper atoms have nuclei that have 29 protons. Each of these atoms also has 29 electrons. Any atom that has 29 protons must be a copper atom. In contrast, knowing just the mass number does not help you identify the element. For example, some copper atom nuclei have 36 neutrons. These copper atoms have a mass number of 65. But zinc atoms that have 30 protons and 35 neutrons also have mass numbers of Chapter 3

16 Atomic Structures Can Be Represented by Symbols Each element has a name, and the same name is given to all atoms of an element. For example, sulfur is composed of sulfur atoms. Recall that each element also has a symbol, and the same symbol is used to represent one of the element s atoms. Thus, S represents a single atom of sulfur, 2S represents two sulfur atoms, and 8S represents eight sulfur atoms. However, chemists write S 8 to indicate that the eight sulfur atoms are joined together and form a molecule of sulfur, as shown in the model in Figure 13. Atomic number and mass number are sometimes written with an element s symbol. The atomic number always appears on the lower left side of the symbol. For example, the symbols for the first five elements are written with atomic numbers as follows: 1H 2He 3Li 4Be 5B Note that these subscript numbers give no new information. They simply indicate the atomic number of a particular element. On the other hand, mass numbers provide information that specifies particular atoms of an element. Mass numbers are written on the upper left side of the symbol. The following are the symbols of stable atoms of the first five elements written with mass numbers: Topic : Atomic Structures SciLinks code: HW H 2 H 3 He 4 He 6 Li 7 Li 9 Be 10 B 11 B Both numbers may be written with the symbol. For example, the most abundant kind of each of the first five elements can be represented by the following symbols: 1 1H 4 2He 7 3Li 9 4Be 11 5B An element may be represented by more than one notation. For example, the following notations represent the different atoms of hydrogen: 1 1H 2 1H 3 1H Hydrogen, H 2 Figure 13 In nature, elemental sulfur exists as eight sulfur atoms joined in a ring, elemental hydrogen exists as a molecule of two hydrogen atoms, and elemental helium exists as single helium atoms. Sulfur, S 8 Helium, He Atoms and Moles 87

17 Figure 14 The two stable isotopes of helium are helium-3 and helium-4. The nucleus of a helium-4 atom is known as an alpha particle. Proton Neutron Proton Neutron Electron cloud Electron cloud Helium-3 Helium-4 isotope an atom that has the same number of protons (atomic number) as other atoms of the same element but has a different number of neutrons (atomic mass) Isotopes of an Element Have the Same Atomic Number All atoms of an element have the same atomic number and the same number of protons. However, atoms do not necessarily have the same number of neutrons. Atoms of the same element that have different numbers of neutrons are called isotopes. The two atoms modeled in Figure 14 are stable isotopes of helium. There are two standard methods of identifying isotopes. One method is to write the mass number with a hyphen after the name of an element. For example, the helium isotope shown on the left in Figure 14 is written helium-3, while the isotope shown on the right is written as helium-4. The second method shows the composition of a nucleus as the isotope s nuclear symbol. Using this method, the notations for the two helium isotopes shown in Figure 14 are written below. 3 2He 4 2He Topic : Atoms and Elements SciLinks code: HW4017 Notice that all isotopes of an element have the same atomic number. However, their atomic masses are not the same because the number of neutrons of the atomic nucleus of each isotope varies. In the case of helium, both isotopes have two protons in their nuclei. However, helium-3 has one neutron, while helium-4 has two neutrons. Table 4 lists the four stable isotopes of lead. The least abundant of these isotopes is lead-204, while the most common is lead-208. Why do all lead atoms have 82 protons and 82 electrons? Table 4 The Stable Isotopes of Lead Name of atom Symbol Number of neutrons Mass number Mass (kg) Abundance (%) Lead-204 Lead-206 Lead-207 Lead Pb Pb Pb Pb Chapter 3

18 SAMPLE PROBLEM B Determining the Number of Particles in Isotopes Calculate the numbers of protons, electrons, and neutrons in oxygen-17 and in oxygen Gather information. The mass numbers for the two isotopes are 17 and Plan your work. An oxygen atom must be electrically neutral. 3 Calculate. atomic number = number of protons = number of electrons = 8 mass number atomic number = number of neutrons For oxygen-17, 17 8 = 9 neutrons For oxygen-18, 18 8 = 10 neutrons 4 Verify your results. The two isotopes have the same numbers of protons and electrons and differ only in their numbers of neutrons. PRACTICE 1 Chlorine has two stable isotopes, chlorine-35 and chlorine-37. The atomic number of chlorine is 17. Calculate the numbers of protons, electrons, and neutrons each isotope has. 2 Calculate the numbers of protons, electrons, and neutrons for each of the following isotopes of calcium: 42 20Ca and 44 20Ca. PRACTICE HINT The only difference between the isotopes of an element is the number of neutrons in the atoms of each isotope. PROBLEM SOLVING SKILL 2 Section Review UNDERSTANDING KEY IDEAS 1. Describe the differences between electrons, protons, and neutrons. 2. How are isotopes of the same element alike? 3. What subatomic particle was discovered with the use of a cathode-ray tube? PRACTICE PROBLEMS 4. Write the symbol for element X, which has 22 electrons and 22 neutrons. 5. Determine the numbers of electrons, protons, and neutrons for each of the following: a Br b Pd c Cs 6. Calculate the atomic number and mass number of an isotope that has 56 electrons and 82 neutrons. CRITICAL THINKING 7. Why must there be an attractive force to explain the existence of stable nuclei? 8. Are hydrogen-3 and helium-3 isotopes of the same element? Explain your answer. Atoms and Moles 89

19 SECTION 3 Electron Configuration KEY TERMS orbital electromagnetic spectrum ground state excited state quantum number Pauli exclusion principle electron configuration aufbau principle Hund s rule Figure 15 According to Rutherford s model of the atom, electrons orbit the nucleus just as planets orbit the sun OBJECTIVES Compare the Rutherford, Bohr, and quantum models of an atom. Explain how the wavelengths of light emitted by an atom provide information about electron energy levels. List the four quantum numbers, and describe their significance. Write the electron configuration of an atom by using the Pauli exclusion principle and the aufbau principle. Atomic Models Soon after the atomic theory was widely accepted by scientists, they began constructing models of atoms. Scientists used the information that they had about atoms to build these models. They knew, for example, that an atom has a densely packed nucleus that is positively charged. This conclusion was the only way to explain the data from Rutherford s gold foil experiments. Building a model helps scientists imagine what may be happening at the microscopic level. For this very same reason, the illustrations in this book provide pictures that are models of chemical compounds to help you understand the relationship between the macroscopic and microscopic worlds. Scientists knew that any model they make may have limitations. A model may even have to be modified or discarded as new information is found. This is exactly what happened to scientists models of the atom. Rutherford s Model Proposed Electron Orbits The experiments of Rutherford s team led to the replacement of the plumpudding model of the atom with a nuclear model of the atom. Rutherford suggested that electrons, like planets orbiting the sun, revolve around the nucleus in circular or elliptical orbits. Figure 15 shows Rutherford s model. Because opposite charges attract, the negatively charged electrons should be pulled into the positively charged nucleus. Because Rutherford s model could not explain why electrons did not crash into the nucleus, his model had to be modified. The Rutherford model of the atom, in turn, was replaced only two years later by a model developed by Niels Bohr, a Danish physicist. The Bohr model, which is shown in Figure 16, describes electrons in terms of their energy levels. 90 Chapter 3

20 Bohr s Model Confines Electrons to Energy Levels According to Bohr s model, electrons can be only certain distances from the nucleus. Each distance corresponds to a certain quantity of energy that an electron can have.an electron that is as close to the nucleus as it can be is in its lowest energy level.the farther an electron is from the nucleus, the higher the energy level that the electron occupies. The difference in energy between two energy levels is known as a quantum of energy. The energy levels in Bohr s model can be compared to the rungs of a ladder. A person can go up and down the ladder only by stepping on the rungs. When standing on the first rung, the person has the lowest potential energy. By climbing to the second rung, the person increases his or her potential energy by a fixed, definite quantity. Because the person cannot stand between the rungs on the ladder, the person s potential energy cannot have a continuous range of values. Instead, the values can be only certain, definite ones. In the same way, Bohr s model states that an electron can be in only one energy level or another, not between energy levels. Bohr also concluded that an electron did not give off energy while in a given energy level. Electrons Act Like Both Particles and Waves Thomson s experiments demonstrated that electrons act like particles that have mass. Although the mass of an electron is extremely small, electrons in a cathode ray still have enough mass to turn a paddle wheel. In 1924, Louis de Broglie pointed out that the behavior of electrons according to Bohr s model was similar to the behavior of waves. For example, scientists knew that any wave confined in space can have only certain frequencies. The frequency of a wave is the number of waves that pass through a given point in one second. De Broglie suggested that electrons could be considered waves confined to the space around a nucleus. As waves, electrons could have only certain frequencies. These frequencies could correspond to the specific energy levels in which electrons are found. Other experiments also supported the wave nature of electrons. Like light waves, electrons can change direction through diffraction. Diffraction refers to the bending of a wave as the wave passes by the edge of an object, such as a crystal. Experiments also showed that electron beams, like waves, can interfere with each other. Figure 17 shows the present-day model of the atom, which takes into account both the particle and wave properties of electrons. According to this model, electrons are located in orbitals, regions around a nucleus that correspond to specific energy levels. Orbitals are regions where electrons are likely to be found. Orbitals are sometimes called electron clouds because they do not have sharp boundaries. When an orbital is drawn, it shows where electrons are most likely to be. Because electrons can be in other places, the orbital has a fuzzy boundary like a cloud. As an analogy to an electron cloud, imagine the spinning blades of a fan. You know that each blade can be found within the spinning image that you see. However, you cannot tell exactly where any one blade is at a particular moment. Figure 16 According to Bohr s model of the atom, electrons travel around the nucleus in specific energy levels. orbital a region in an atom where there is a high probability of finding electrons Figure 17 According to the current model of the atom, electrons are found in orbitals. Atoms and Moles 91

21 Topic : Electromagnetic Spectrum SciLinks code: HW4048 electromagnetic spectrum all of the frequencies or wavelengths of electromagnetic radiation Figure 18 The electromagnetic spectrum is composed of light that has a broad range of wavelengths. Our eyes can detect only the visible spectrum. Electrons and Light By 1900, scientists knew that light could be thought of as moving waves that have given frequencies, speeds, and wavelengths. In empty space, light waves travel at m/s.at this speed, light waves take only 500 s to travel the 150 million kilometers between the sun and Earth. The wavelength is the distance between two consecutive peaks or troughs of a wave. The distance of a wavelength is usually measured in meters. The wavelength of light can vary from 10 5 m to less than m. This broad range of wavelengths makes up the electromagnetic spectrum, which is shown in Figure 18. Notice in Figure 18 that our eyes are sensitive to only a small portion of the electromagnetic spectrum. This sensitivity ranges from 700 nm, which is about the value of wavelengths of red light, to 400 nm, which is about the value of wavelengths of violet light. In 1905, Albert Einstein proposed that light also has some properties of particles. His theory would explain a phenomenon known as the photoelectric effect. This effect happens when light strikes a metal and electrons are released. What confused scientists was the observation that for a given metal, no electrons were emitted if the light s frequency was below a certain value, no matter how long the light was on. Yet if light were just a wave, then any frequency eventually should supply enough energy to remove an electron from the metal. Einstein proposed that light has the properties of both waves and particles. According to Einstein, light can be described as a stream of particles, the energy of which is determined by the light s frequency. To remove an electron, a particle of light has to have at least a minimum energy and therefore a minimum frequency. Visible spectrum Violet Blue Green Yellow Orange Red 400 nm 500 nm 600 nm 700 nm rays X rays Ultraviolet Infrared Microwaves Radio waves 1 pm 10 pm 0.1 nm 1 nm 10 nm 0.1 m 1 m 10 m 0.1 mm 1 mm 10 mm 0.1 m 1 m 10 m 0.1 km 1 km 10 km Wavelength Frequency Hz Hz Hz Hz Hz Hz Hz Hz Hz Hz 10 9 Hz 100 MHz 10 MHz 1 MHz 100 khz 92 Chapter 3 Electromagnetic spectrum

22 Red light Low frequency Long wavelength Figure 19 The frequency and wavelength of a wave are inversely related. As frequency increases, wavelength decreases. Violet light High frequency Short wavelength Light Is an Electromagnetic Wave When passed through a glass prism, sunlight produces the visible spectrum all of the colors of light that the human eye can see. You can see from Figure 18 on the previous page that the visible spectrum is only a tiny portion of the electromagnetic spectrum. The electromagnetic spectrum also includes X rays, ultraviolet and infrared light, microwaves, and radio waves. Each of these electromagnetic waves is referred to as light, although we cannot see these wavelengths. Figure 19 shows the frequency and wavelength of two regions of the spectrum that we see: red and violet lights. If you compare red and violet lights, you will notice that red light has a low frequency and a long wavelength. But violet light has a high frequency and a short wavelength. The frequency and wavelength of a wave are inversely related. Light Emission When a high-voltage current is passed through a tube of hydrogen gas at low pressure, lavender-colored light is seen. When this light passes through a prism, you can see that the light is made of only a few colors. This spectrum of a few colors is called a line-emission spectrum. Experiments with other gaseous elements show that each element has a line-emission spectrum that is made of a different pattern of colors. In 1913, Bohr showed that hydrogen s line-emission spectrum could be explained by assuming that the hydrogen atom s electron can be in any one of a number of distinct energy levels. The electron can move from a low energy level to a high energy level by absorbing energy. Electrons at a higher energy level are unstable and can move to a lower energy level by releasing energy. This energy is released as light that has a specific wavelength. Each different move from a particular energy level to a lower energy level will release light of a different wavelength. Bohr developed an equation to calculate all of the possible energies of the electron in a hydrogen atom. His values agreed with those calculated from the wavelengths observed in hydrogen s line-emission spectrum. In fact, his values matched with the experimental values so well that his atomic model that is described earlier was quickly accepted. Topic : Light and Color SciLinks code: HW Topic : Producing Light SciLinks code: HW4099 Atoms and Moles 93

23 ground state the lowest energy state of a quantized system excited state a state in which an atom has more energy than it does at its ground state Topic : Energy Levels SciLinks code: HW4051 Light Provides Information About Electrons Normally, if an electron is in a state of lowest possible energy, it is in a ground state. If an electron gains energy, it moves to an excited state. An electron in an excited state will release a specific quantity of energy as it quickly falls back to its ground state. This energy is emitted as certain wavelengths of light, which give each element a unique line-emission spectrum. Figure 20 shows the wavelengths of light in a line-emission spectrum for hydrogen, through which a high-voltage current was passed. The highvoltage current may supply enough energy to move an electron from its ground state, which is represented by n = 1 in Figure 20, to a higher excited state for an electron in a hydrogen atom, represented by n > 1. Eventually, the electron will lose energy and return to a lower energy level. For example, the electron may fall from the n = 7 energy level to the n = 3 energy level. Notice in Figure 20 that when this drop happens, the electron emits a wavelength of infrared light. An electron in the n = 6 energy level can also fall to the n = 2 energy level. In this case, the electron emits a violet light, which has a shorter wavelength than infrared light does. Figure 20 An electron in a hydrogen atom can move between only certain energy states, shown as n = 1 to n = 7. In dropping from a higher energy state to a lower energy state, an electron emits a characteristic wavelength of light. Energy Infrared wavelengths n = n = 7 n = 6 n = 5 n = 4 n = 3 n = 2 Wavelength (nm) Ultraviolet wavelengths n = 1 94 Chapter 3

24 Table 5 Quantum Numbers of the First 30 Atomic Orbitals n l m Orbital name Number of orbitals s s , 0, 1 2p s , 0, 1 3p , 1, 0, 1, 2 3d s , 0, 1 4p , 1, 0, 1, 2 4d , 2, 1, 0, 1, 2, 3 4f 7 Quantum Numbers The present-day model of the atom, in which electrons are located in orbitals, is also known as the quantum model. According to this model, electrons within an energy level are located in orbitals, regions of high probability for finding a particular electron. However, the model does not explain how the electrons move about the nucleus to create these regions. To define the region in which electrons can be found, scientists have assigned four quantum numbers to each electron. Table 5 lists the quantum numbers for the first 30 atomic orbitals. The principal quantum number, symbolized by n, indicates the main energy level occupied by the electron. Values of n are positive integers, such as 1, 2, 3, and 4. As n increases, the electron s distance from the nucleus and the electron s energy increases. The main energy levels can be divided into sublevels. These sublevels are represented by the angular momentum quantum number, l. This quantum number indicates the shape or type of orbital that corresponds to a particular sublevel. Chemists use a letter code for this quantum number. A quantum number l = 0 corresponds to an s orbital, l = 1 to a p orbital, l = 2 to a d orbital, and l = 3 to an f orbital. For example, an orbital with n = 3 and l = 1 is called a 3p orbital, and an electron occupying that orbital is called a 3p electron. The magnetic quantum number, symbolized by m, is a subset of the l quantum number. It also indicates the numbers and orientations of orbitals around the nucleus. The value of m takes whole-number values, depending on the value of l. The number of orbitals includes one s orbital, three p orbitals, five d orbitals, and seven f orbitals. The spin quantum number, symbolized by or 1 2 ( or ), indicates the orientation of an electron s magnetic field relative to an outside magnetic field. A single orbital can hold a maximum of two electrons, which must have opposite spins. quantum number a number that specifies the properties of electrons Atoms and Moles 95

25 Pauli exclusion principle the principle that states that two particles of a certain class cannot be in the exact same energy state electron configuration the arrangement of electrons in an atom Electron Configurations Figure 21 shows the shapes and orientations of the s, p, and d orbitals. Each orbital that is shown can hold a maximum of two electrons. The discovery that two, but no more than two, electrons can occupy a single orbital was made in 1925 by the German chemist Wolfgang Pauli. This rule is known as the Pauli exclusion principle. Another way of stating the Pauli exclusion principle is that no two electrons in the same atom can have the same four quantum numbers.the two electrons can have the same value of n by being in the same main energy level. These two electrons can also have the same value of l by being in orbitals that have the same shape. And, these two electrons may have the same value of m by being in the same orbital. But these two electrons cannot have the same spin quantum number. If one electron has the value of + 1, then the other electron must have the value of The arrangement of electrons in an atom is usually shown by writing an electron configuration. Like all systems in nature, electrons in atoms tend to assume arrangements that have the lowest possible energies. An electron configuration of an atom shows the lowest-energy arrangement of the electrons for the element. z y z y x x z p x orbital x d x 2 y 2 orbital y z y z y z y x d xz orbital x x x z s orbital p y orbital y d xy orbital y Figure 21 a The s orbital is spherically shaped. There is one s orbital for each value n = 1, 2, 3 of the principal number. z y x z y x d z 2 orbital x p z orbital z d yz orbital b For each of the values n = 2, 3, 4, there are three p orbitals. All are dumbbell shaped, but they differ in orientation. c For each of the values n = 3, 4, 5, there are five d orbitals. Four of the five have similar shapes, but differ in orientation. 96 Chapter 3

26 An Electron Occupies the Lowest Energy Level Available The Pauli exclusion principle is one rule to help you write an electron configuration for an atom. Another rule is the aufbau principle. Aufbau is the German word for building up. The aufbau principle states that electrons fill orbitals that have the lowest energy first. Recall that the smaller the principal quantum number, the lower the energy. But within an energy level, the smaller the l quantum number, the lower the energy. Recall that chemists use letters to represent the l quantum number. So, the order in which the orbitals are filled matches the order of energies, which starts out as follows: 1s < 2s < 2p < 3s < 3p After this point, the order is less obvious. Figure 22 shows that the energy of the 3d orbitals is slightly higher than the energy of the 4s orbitals. As a result, the order in which the orbitals are filled is as follows: 1s < 2s < 2p < 3s < 3p < 4s < 3d Additional irregularities occur at higher energy levels. Can you determine which orbitals electrons of a carbon atom occupy? Two electrons occupy the 1s orbital, two electrons occupy the 2s orbital, and two electrons occupy the 2p orbitals. Now try the same exercise for titanium. Two electrons occupy the 1s orbital, two electrons occupy the 2s orbital, six electrons occupy the 2p orbitals, two electrons occupy the 3s orbital, six electrons occupy the 3p orbitals, two electrons occupy the 3d orbitals, and two electrons occupy the 4s orbital. aufbau principle the principle that states that the structure of each successive element is obtained by adding one proton to the nucleus of the atom and one electron to the lowest-energy orbital that is available Energy n = 4 n = 3 4f 4d 4p 3d 4s 3p 3s Figure 22 This diagram illustrates how the energy of orbitals can overlap such that 4s fills before 3d. n = 2 2p 2s n = 1 1s Atoms and Moles 97

27 An Electron Configuration Is a Shorthand Notation Based on the quantum model of the atom, the arrangement of the electrons around the nucleus can be shown by the nucleus s electron configuration. For example, sulfur has sixteen electrons. Its electron configuration is written as 1s 2 2s 2 2p 6 3s 2 3p 4. This line of symbols tells us about these sixteen electrons. Two electrons are in the 1s orbital, two electrons are in the 2s orbital, six electrons are in the 2p orbitals, two electrons are in the 3s orbital, and four electrons are in the 3p orbitals. Each element s configuration builds on the previous elements configurations. To save space, one can write this configuration by using a configuration of a noble gas. The noble gas electron configurations that are often used are the configurations of neon, argon, krypton, and xenon. The neon atom s configuration is 1s 2 2s 2 2p 6, so the electron configuration of sulfur is written as shown below. [Ne] 3s 2 3p 4 Hund s rule the rule that states that for an atom in the ground state, the number of unpaired electrons is the maximum possible and these unpaired electrons have the same spin Does an electron enter the first 3p orbital to pair with a single electron that is already there? Or does the electron fill another 3p orbital? According to Hund s rule, the second answer is correct. Hund s rule states that orbitals of the same n and l quantum numbers are each occupied by one electron before any pairing occurs. For example, sulfur s configuration is shown by the orbital diagram below. Electrons are represented by arrows. Note that an electron fills another orbital before the electron occupies an orbital that occupied. 1s 2s 2p 3s 3p SAMPLE PROBLEM C PRACTICE HINT Remember that an s orbital holds 2 electrons, three p orbitals hold 6 electrons, and five d orbitals hold 10 electrons. Writing Electron Configurations Write the electron configuration for an atom whose atomic number is Gather information. The atomic number of the element is Plan your work. The atomic number represents the number of protons in an atom. The number of protons must equal the number of electrons in an atom. Write the electron configuration for that number of electrons by following the Pauli exclusion principle and the aufbau principle. A noble gas configuration can be used to write this configuration. 98 Chapter 3

28 3 Calculate. atomic number = number of protons = number of electrons = 20 According to the aufbau principle, the order of orbital filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. The electron configuration for an atom of this element is written as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 This electron configuration can be abbreviated as follows: [Ar]4s 2 4 Verify your results. The sum of the superscripts is ( ) = 20. Therefore, all 20 electrons are included in the electron configuration. PRACTICE 1 Write the electron configuration for an atom of an element whose atomic number is 8. 2 Write the electron configuration for an atom that has 17 electrons. PROBLEM SOLVING SKILL 3 Section Review UNDERSTANDING KEY IDEAS 1. How does Bohr s model of the atom differ from Rutherford s? 2. What happens when an electron returns to its ground state from its excited state? 3. What does n represent in the quantum model of electrons in atoms? PRACTICE PROBLEMS 4. What is the atomic number of an element whose atom has the following electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2? 5. Write the electron configuration for an atom that has 13 electrons. 6. Write the electron configuration for an atom that has 33 electrons. 7. How many orbitals are completely filled in an atom whose electron configuration is 1s 2 2s 2 2p 6 3s 1? CRITICAL THINKING 8. Use the Pauli exclusion principle or the aufbau principle to explain why the following electron configurations are incorrect: a. 1s 2 2s 3 2p 6 3s 1 b. 1s 2 2s 2 2p 5 3s 1 9. Why is a shorter wavelength of light emitted when an electron falls from n = 4 to n = 1 than when an electron falls from n = 2 to n = 1? 10. Calculate the maximum number of electrons that can occupy the third principal energy level. 11. Why do electrons fill the 4s orbital before they start to occupy the 3d orbital? Atoms and Moles 99

29 SECTION 4 Counting Atoms KEY TERMS atomic mass mole molar mass Avogadro s number OBJECTIVES Compare the quantities and units for atomic mass with those for molar mass. Define mole, and explain why this unit is used to count atoms. Calculate either mass with molar mass or number with Avogadro s number given an amount in moles. Atomic Mass You would not expect something as small as an atom to have much mass. For example, copper atoms have an average mass of only kg. Each penny in Figure 23 has an average mass of kg and contains copper. How many copper atoms are there in one penny? Assuming that a penny is pure copper, you can find the number of copper atoms by dividing the mass of the penny by the average mass of a single copper atom or by using the following conversion factor: 1 atom Cu/ kg atom Cu kg = Cu atoms kg atomic mass the mass of an atom expressed in atomic mass units Figure 23 These pennies are made mostly of copper atoms. Each copper atom has an average mass of kg. Masses of Atoms Are Expressed in Atomic Mass Units Obviously, atoms are so small that the gram is not a very convenient unit for expressing their masses. Even the picogram (10 12 g) is not very useful. A special mass unit is used to express atomic mass. This unit has two names the atomic mass unit (amu) and the Dalton (Da). In this book, atomic mass unit will be used. But how can you tell what the atomic mass of a specific atom is? When the atomic mass unit was first set up, an atom s mass number was supposed to be the same as the atom s mass in atomic mass units. Mass number and atomic mass units would be the same because a proton and a neutron each have a mass of about 1.0 amu. For example, a copper-63 atom has an atomic mass of A copper-65 atom has an atomic mass of (The slight differences from exact values will be discussed in later chapters.) Another way to determine atomic mass is to check a periodic table, such as the one on the inside cover of this book. The mass shown is an average of the atomic masses of the naturally occurring isotopes. For this reason, copper is listed as instead of or Chapter 3

30 Introduction to the Mole Most samples of elements have great numbers of atoms. To make working with these numbers easier, chemists created a new unit called the mole (mol). A mole is defined as the number of atoms in exactly 12 grams of carbon-12. The mole is the SI unit for the amount of a substance. Chemists use the mole as a counting unit, just as you use the dozen as a counting unit. Instead of asking for 12 eggs, you ask for 1 dozen eggs. Similarly, chemists refer to 1 mol of carbon or 2 mol of iron. To convert between moles and grams, chemists use the molar mass of a substance. The molar mass of an element is the mass in grams of one mole of the element. Molar mass has the unit grams per mol (g/mol).the mass in grams of 1 mol of an element is numerically equal to the element s atomic mass from the periodic table in atomic mass units. For example, the atomic mass of copper to two decimal places is amu. Therefore, the molar mass of copper is g/mol. Skills Toolkit 1 shows how to convert between moles and mass in grams using molar mass. Scientists have also determined the number of particles present in 1 mol of a substance, called Avogadro s number. One mole of pure substance contains particles. To get some idea of how large Avogadro s number is, imagine that every living person on Earth (about 6 billion people) started counting the number of atoms of 1 mol C. If each person counted nonstop at a rate of one atom per second, it would take over 3 million years to count every atom. Avogadro s number may be used to count any kind of particle, including atoms and molecules. For calculations in this book, Avogadro s number will be rounded to particles per mole. Skills Toolkit 2 shows how to use Avogadro s number to convert between amount in moles and the number of particles. mole the SI base unit used to measure the amount of a substance whose number of particles is the same as the number of atoms in 12 g of carbon-12 molar mass the mass in grams of 1 mol of a substance Avogadro s number , the number of atoms or molecules in 1 mol SKILLS 1 Determining the Mass from the Amount in Moles amount mol g 1 mol use molar mass 1 mol g mass g Atoms and Moles 101

31 SAMPLE PROBLEM D Converting from Amount in Moles to Mass Determine the mass in grams of 3.50 mol of copper. 1 Gather information. amount of Cu = 3.50 mol mass of Cu =? g Cu molar mass of Cu = g 2 Plan your work. First, make a set-up that shows what is given and what is desired mol Cu? =? g Cu PRACTICE HINT For elements and compounds, the mass will always be a number that is greater than the number of moles. Use a conversion factor that has g Cu in the numerator and mol Cu in the denominator.? g Cu 3.50 mol Cu =? g Cu 1 molcu 3 Calculate. The correct conversion factor is the molar mass of Cu, g/mol. Place the molar mass in the equation, and calculate the answer. Use the periodic table in this book to find mass numbers of elements mol Cu g Cu = 222 g Cu 1 mol Cu 4 Verify your results. To verify that the answer of 222 g is correct, find the number of moles of 222 g of copper. 1 mol Cu 222 g of Cu = 3.49 mol Cu g Cu The amount of 3.49 mol is close to the 3.50 mol, so the answer of 222 g is reasonable. PROBLEM SOLVING SKILL PRACTICE 1 What is the mass in grams of 1.00 mol of uranium? 2 What is the mass in grams of mol of uranium? 3 Calculate the number of moles of g of hydrogen atoms. What is the mass in grams of mol of hydrogen atoms? 4 Calculate the mass in grams of mol of lead. Calculate the number of moles of g of lead. 102 Chapter 3

32 Determining the Number of Atoms from the Amount in Moles amount SKILLS 2 mol x 10 atoms 1 mol use Avogadro's number 1 mol x 10 atoms number of atoms atoms SAMPLE PROBLEM E Converting from Amount in Moles to Number of Atoms Determine the number of atoms in 0.30 mol of fluorine atoms. 1 Gather information. amount of F = 0.30 mol number of atoms of F =? 2 Plan your work. To determine the number of atoms, select the conversion factor that will take you from the amount in moles to the number of atoms. amount (mol) atoms/mol = number of atoms 3 Calculate F atoms 0.30 mol F = F atoms 1 mol F PRACTICE HINT Make sure to select the correct conversion factor so that units cancel to give the unit required in the answer. 4 Verify your results. The answer has units that are requested in the problem. The answer is also less than atoms, which makes sense because you started with less than 1 mol. PRACTICE 1 How many atoms are in 0.70 mol of iron? 2 How many moles of silver are represented by atoms? 3 How many moles of osmium are represented by atoms? PROBLEM SOLVING SKILL Atoms and Moles 103

33 Figure 24 Carbon, which composes diamond, is the basis for the atomic mass scale that is used today. Chemists and Physicists Agree on a Standard The atomic mass unit has been defined in a number of different ways over the years. Originally, atomic masses expressed the ratio of the mass of an atom to the mass of a hydrogen atom. Using hydrogen as the standard turned out to be inconvenient because hydrogen does not react with many elements. Early chemists determined atomic masses by comparing how much of one element reacted with another element. Because oxygen combines with almost all other elements, oxygen became the standard of comparison. The atomic mass of oxygen was defined as exactly 16, and the atomic masses of the other elements were based on this standard. But this choice also led to difficulties. Oxygen exists as three isotopes. Physicists based their atomic masses on the assignment of as the mass of the most common oxygen isotope. Chemists, on the other hand, decided that should be the average mass of all oxygen isotopes, weighted according to the abundance of each isotope. So, to a physicist, the atomic mass of fluorine was , but to a chemist, it was Finally, in 1962, a conference of chemists and physicists agreed on a scale based on an isotope of carbon. Carbon is shown in Figure 24. Used by all scientists today, this scale defines the atomic mass unit as exactly one-twelfth of the mass of one carbon-12 atom. As a result, one atomic mass unit is equal to kg. The mass of an atom is indeed quite small. 4 Section Review UNDERSTANDING KEY IDEAS 1. What is atomic mass? 2. What is the SI unit for the amount of a substance that contains as many particles as there are atoms in exactly 12 grams of carbon-12? 3. Which atom is used today as the standard for the atomic mass scale? 4. What unit is used for molar mass? 5. How many particles are present in 1 mol of a pure substance? PRACTICE PROBLEMS 6. Convert atoms of silicon to moles of silicon. 7. How many atoms are present in 4.0 mol of sodium? 8. How many moles are represented by 118 g of cobalt? Cobalt has an atomic mass of amu. 9. How many moles are represented by 250 g of platinum? 10. Convert 0.20 mol of boron into grams of boron. How many atoms are present? CRITICAL THINKING 11. What is the molar mass of an element? 12. How is the mass in grams of an element converted to amount in moles? 13. How is the mass in grams of an element converted to number of atoms? 104 Chapter 3

34 BERYLLIUM 4 Be Where Is Be? Earth s crust: 0.005% by mass Element Spotlight Beryllium [He]2s 2 Beryllium: An Uncommon Element Although it is an uncommon element, beryllium has a number of properties that make it very useful. Beryllium has a relatively high melting point (1278 C) and is an excellent conductor of energy as heat and electrical energy. Beryllium transmits X rays extremely well and is therefore used to make windows for X-ray devices. All compounds of beryllium are toxic to humans. People who experience prolonged exposure to beryllium dust may contract berylliosis, a disease that can lead to severe lung damage and even death. Industrial Uses The addition of 2% beryllium to copper forms an alloy that is six times stronger than copper is. This alloy is used for nonsparking tools, critical moving parts in jet engines, and components in precision equipment. Beryllium is used in nuclear reactors as a neutron reflector and as an alloy with the fuel elements. Real-World Connection Emerald and aquamarine are precious forms of the mineral beryl, Be 3 Al 2 (SiO 3 ) 6. Crystals of pure beryllium look very different from the combined form of beryllium in an emerald. A Brief History : F. Wöhler of Germany gives beryllium its name after he and W. Bussy of France simultaneously isolate the pure metal : M. G. Corson of the United States discovers that beryllium can be used to age-harden copper-nickel alloys. 1798: R. J. Haüy, a French mineralogist, observes that emeralds and beryl have the same optical properties and therefore the same chemical composition. 1898: P. Lebeau discovers a method of extracting highpurity beryllium by using an electrolytic process. 1942: A Ra-Be source provides the neutrons for Fermi s studies. These studies lead to the construction of a nuclear reactor. Questions 1. Research how the beryllium and copper alloy is made and what types of equipment are made of this alloy. 2. Research how beryllium is used in nuclear reactors. Topic : Beryllium SciLinks code: HW Research berylliosis and use the information to make a medical information brochure. Be sure to include symptoms, causes, and risk factors in your report. Atoms and Moles 105

35 3 CHAPTER HIGHLIGHTS KEY IDEAS SECTION ONE Substances Are Made of Atoms Three laws support the existence of atoms: the law of definite proportions, the law of conservation of mass, and the law of multiple proportions. Dalton s atomic theory contains five basic principles, some of which have been modified. KEY TERMS law of definite proportions law of conservation of mass law of multiple proportions SECTION TWO Structure of Atoms Protons, particles that have a positive charge, and neutrons, particles that have a neutral charge, make up the nuclei of most atoms. Electrons, particles that have a negative charge and very little mass, occupy the region around the nucleus. The atomic number of an atom is the number of protons the atom has. The mass number of an atom is the number of protons plus the number of neutrons. Isotopes are atoms that have the same number of protons but different numbers of neutrons. SECTION THREE Electron Configuration The quantum model describes the probability of locating an electron at any place. Each electron is assigned four quantum numbers that describe it. No two electrons of an atom can have the same four quantum numbers. The electron configuration of an atom reveals the number of electrons an atom has. SECTION FOUR Counting Atoms The masses of atoms are expressed in atomic mass units (amu). The mass of an atom of the carbon-12 isotope is defined as exactly 12 atomic mass units. The mole is the SI unit for the amount of a substance that contains as many particles as there are atoms in exactly 12 grams of carbon-12. Avogadro s number, particles per mole, is the number of particles in a mole. electron nucleus proton neutron atomic number mass number isotope orbital electromagnetic spectrum ground state excited state quantum number Pauli exclusion principle electron configuration aufbau principle Hund s rule atomic mass mole molar mass Avogadro s number KEY SKILLS Determining the Number of Particles in an Atom Sample Problem A p. 86 Determining the Number of Particles in Isotopes Sample Problem B p. 89 Writing Electron Configurations Sample Problem C p. 98 Converting Amount in Moles to Mass Skills Toolkit 1 p. 101 Sample Problem D p. 102 Converting Amount in Moles to Number of Atoms Skills Toolkit 2 p. 103 Sample Problem E p Chapter 3

36 CHAPTER REVIEW 3 USING KEY TERMS 1. Define isotope. 2. What are neutrons? 3. State the Pauli exclusion principle. 4. What is a cathode? 5. Define mass number. 6. What is a line-emission spectrum? 7. Define ground state. 8. Define mole. 9. State the law of definite proportions. 10. What is an orbital? 11. What is an electron configuration? UNDERSTANDING KEY IDEAS Substances Are Made of Atoms 12. What law is illustrated by the fact that ice, water, and steam consist of 88.8% oxygen and 11.2% hydrogen by mass? 13. What law is shown by the fact that the mass of carbon dioxide, which forms as a product of a reaction between oxygen and carbon, equals the combined masses of the carbon and oxygen that reacted? 14. Of the five parts of Dalton s atomic theory, which one(s) have been modified? Structure of Atoms 15. How were atomic models developed given that no one had seen an atom? 16. Why are atomic numbers always whole numbers? 17. If all protons have positive charges, how can any atomic nucleus be stable? 18. What observation did Thomson make to suggest that an electron has a negative electric charge? Electron Configuration 19. How do you use the aufbau principle when you create an electron configuration? 20. Explain what is required to move an electron from the ground state to an excited state. 21. Why can a p sublevel hold six electrons while the s sublevel can hold no more than two electrons? 22. What do electrons and light have in common? 23. How are the frequency and wavelength of light related? 24. Why does an electron occupy the 4s orbital before the 3d orbital? 25. The element sulfur has an electron configuration of 1s 2 2s 2 2p 6 3s 2 3p 4. a. What does the superscript 6 refer to? b. What does the letter s refer to? c. What does the coefficient 3 refer to? Counting Atoms 26. What is a mole? How is a mole related to Avogadro s number? 27. What significance does carbon-12 have in terms of atomic mass? 28. If the mass of a gold atom is amu, what is the atom s molar mass? 29. What advantage is gained by using the mole as a unit when working with atoms? Atoms and Moles 107

37 PRACTICE PROBLEMS Sample Problem A Determining the Number of Particles in an Atom 30. Calculate the number of neutrons of the atom whose atomic number is 42 and whose mass number is How many electrons are present in an atom of mercury whose atomic number is 80 and whose mass number is 201? 32. Calculate the number of protons of the atom whose mass number is 19 and whose number of neutrons is Calculate the number of electrons of the atom whose mass number is 75 and whose number of neutrons is 42. Sample Problem B Determining the Number of Particles in Isotopes 34. Write nuclear symbols for isotopes of uranium that have the following numbers of neutrons. The atomic number of uranium is 92. a. 142 neutrons b. 143 neutrons c. 146 neutrons 35. Copy and complete the following table concerning the three isotopes of silicon, whose atomic number is 14. Isotope Number Number Number of protons of electrons of neutrons Si-28 Si-29 Si-30 PROBLEM SOLVING SKILL 36. Write the symbol for two isotopes of carbon. Both isotopes have six protons. One isotope has six neutrons, while the other has seven neutrons. 37. All barium atoms have 56 protons. One isotope of barium has 74 neutrons, and another isotope has 81 neutrons. Write the symbols for these two isotopes of barium. Sample Problem C Writing Electron Configurations 38. Write the electron configuration for nickel, whose atomic number is 28. Remember that the 4s orbital has lower energy than the 3d orbital does and that the d sublevel can hold a maximum of 10 electrons. 39. Write the electron configuration of germanium whose atomic number is How many orbitals are completely filled in an atom that has 12 electrons? The electron configuration is 1s 2 2s 2 2p 6 3s How many orbitals are completely filled in an atom of an element whose atomic number is 18? Sample Problem D Moles to Mass Converting Amount in 42. How many moles are represented by each of the following. a g Na which has an atomic mass of amu b. 150 g S which has an atomic mass of amu c g Ni which has an atomic mass of amu 43. Determine the mass in grams represented by 2.50 mol tellurium. 44. What is the mass in grams of mol of hydrogen atoms? Sample Problem E Converting Amount in Moles to Number of Atoms 45. Calculate the number of atoms in 2.0 g of hydrogen atoms. The atomic mass of hydrogen is 1.01 amu. 46. Calculate the number of atoms present in each of the following: a. 2 mol Fe b g Ca, which has an atomic mass of amu c. 4.5 mol of boron Chapter 3

38 47. How many mol of potassium are represented by potassium atoms? MIXED REVIEW 48. In the diagram below, indicate which subatomic particles would be found in areas a and b. a 49. What mass of silver, Ag, which has an atomic mass of amu, contains the same number of atoms contained in 10.0 g of boron, B, which has an atomic mass of amu? 50. Hydrogen s only electron occupies the 1s orbital but can be excited to a 4p orbital. List all of the orbitals that this electron can occupy as it falls. 51. What is the electron configuration of zinc? 52. Identify the scientists who proposed each of the models illustrated below. a. c. b. d. b 54. How did the results of the gold foil experiment lead Rutherford to recognize the existence of atomic nuclei? 55. Explain why atoms are neutral. 56. Explain Coulomb s law. 57. Determine the mass in kilograms of 5.50 mol of iron, Fe. 58. What is Avogadro s number? 59. How many moles are present in 11 g of silicon? how many atoms? 60. Suppose an atom has a mass of 11 amu and has five electrons. What is this atom s atomic number? 61. Explain why different atoms of the same element always have the same atomic number but can have different mass numbers. 62. What does an element s molar mass tell you about the element? 63. A pure gold bar is made of mol of gold. What is the mass of the bar in grams? 64. Write the electron configuration of phosphorus. 65. What are the charges of an electron, a proton, and a neutron? 66. An advertising sign gives off red and green light. a. Which light has higher energy? b. One of the colors has a wavelength of 680 nm and the other has a wavelength of 500. Which color has which wavelength? 67. Can a stable atom have an orbital which has three electrons? Explain your answer. CRITICAL THINKING 53. How many atoms are in 0.75 moles of neptunium? 68. Predict what Rutherford might have observed if he had bombarded copper metal instead of gold metal with alpha particles. The atomic numbers of copper and gold are 29 and 79, respectively. Atoms and Moles 109

39 69. Identify the law that explains why a water molecule in a raindrop falling on Phoenix, Arizona, and a water molecule in the Nile River in Egypt are both made of two hydrogen atoms for every oxygen atom. 70. Which of Dalton s principles is contradicted by a doctor using radioisotopes to trace chemicals in the body? 71. For hundreds of years, alchemists searched for ways to turn various metals into gold. How would the structure of an atom of Hg (mercury) have to be changed for the atom to become an atom of Au (gold)? 72. How are quantum numbers like an address? How are they different from an address? 73. Which has more atoms: 3.0 g of iron, Fe, or 2.0 grams of sulfur, S? 74. Predict which isotope of nitrogen is more commonly found, nitrogen-14 or nitrogen Suppose you have only 1.9 g of sulfur for an experiment and you must do three trials using mol of S each time. Do you have enough sulfur? 76. How many orbitals in an atom can have the following designation? a. 4p b. 7s c. 5d 77. Explain why that if n = 2, l cannot be Write the electron configuration of tin. 79. Many elements exist as polyatomic molecules. Use atomic masses to calculate the molecular masses of the following: a. O 2 b. P 4 c. S What do the electron configurations of neon, argon, krypton, xenon, and radon have in common? 81. The magnetic properties of an element depend on the number of unpaired electrons it has. Explain why iron, Fe, is highly magnetic but neon, Ne, is not. 82. Answer the following regarding electron configurations of atoms in the fourth period of the periodic table. a. Which orbitals are filled by transition metals? b. Which orbitals are filled by nonmetals? ALTERNATIVE ASSESSMENT 83. So-called neon signs actually contain a variety of gases. Research the different substances used for these signs. Design your own sign on paper, and identify which gases you would use to achieve the desired color scheme. 84. Research several elements whose symbols are inconsistent with their English names. Some examples include silver, Ag; gold, Au; and mercury, Hg. Compare the origin of these names with the origin of the symbols. 85. Research the development of the scanning tunneling microscope, which can be used to make images of atoms. Find out what information about the structure of atoms these microscopes have provided. 86. Select one of the essential elements. Check your school library or the Internet for details about the role of each element in the human body and for any guidelines and recommendations about the element. CONCEPT MAPPING 87. Use the following terms to create a concept map: proton, atomic number, atomic theory, orbital, and electron. 110 Chapter 3

40 FOCUS ON GRAPHING Study the graph below, and answer the questions that follow. For help in interpreting graphs, see Appendix B, Study Skills for Chemistry. 88. What represents the ground state in this diagram? 89. Which energy-level changes can be detected by the unaided eye? 90. Does infrared light have more energy than ultraviolet light? Why or why not? 91. Which energy levels represent a hydrogen electron in an excited state? 92. What does the energy level labeled n = represent? 93. If an electron is beyond the n = level, is the electron a part of the hydrogen atom? Energy Infrared wavelengths n = n = 7 n = 6 n = 5 n = 4 n = 3 n = 2 Wavelength (nm) Ultraviolet wavelengths n = 1 TECHNOLOGY AND LEARNING 94. Graphing Calculator Calculate Numbers of Protons, Electrons, and Neutrons. A graphing calculator can run a program that calculates the numbers of protons, electrons, and neutrons given the atomic mass and numbers for an atom. For example, given a calcium-40 atom, you will calculate the numbers of protons, electrons, and neutrons in the atom. Go to Appendix C. If you are using a TI-83 Plus, you can download the program NUMBER and data and can run the application as directed. If you are using another calculator, your teacher will provide you with keystrokes and data sets to use. After you have run the program, answer the questions below. a. Which element has the most protons? b. How many neutrons does mercury-201 have? c. Carbon-12 and carbon-14 have the same atomic number. Do they have the same number of neutrons? Why or why not? Atoms and Moles 111

41 3 STANDARDIZED TEST PREP UNDERSTANDING CONCEPTS Directions (1 3): For each question, write on a separate sheet of paper the letter of the correct answer Which of the following represents an electron configuration of a calcium atom, whose atomic number is 20? A. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 B. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 3 C. 1s 2 2s 2 2p 6 3s 1 3p 6 4s 2 3d 1 D. 1s 2 2s 2 2p 6 3s 2 3d 8 Which of these is always equal to the number of protons in an atom? F. the mass number G. the number of isotopes H. the number of neutrons I. the number of electrons Which of these events occurs when an electron in an excited state returns to its ground state? A. Light energy is emitted. B. Energy is absorbed by the atom. C. The atom undergoes spontaneous decay. D. The charge increases because an electron is added. Directions (4 6): For each question, write a short response What is the electron configuration of bromine, whose atomic number is 35? Electrons do not always act like particles. What electron behavior did de Broglie observe, and what evidence did he use to support his ideas? Only materials with unpaired electrons can exhibit magnetic properties. Can the element xenon be highly magnetic? Explain. READING SKILLS Directions (7 9): Read the passage below. Then answer the questions Although there is no detector that allows us to see the inside of an atom, scientists infer its structure from the properties of its components. Rutherford s model shows electrons orbiting the nucleus like planets around the sun. In Bohr s model the electrons travel around the nucleus in specific energy levels. According to the current model, electron orbitals do not have sharp boundaries and the electrons are portrayed as a cloud. The model of the atom has changed over time because F. earlier models were proven to be wrong G. electrons do not revolve around the nucleus H. as new properties of atoms were discovered, models had to be revised to account for those properties I. new particles were discovered, so the model had to be changed to explain how they could exist Why do scientists need models as opposed to directly observing electrons? A. Models can be changed. B. There is no technology that allows direct observation of electrons. C. The charges on the electrons and protons interfere with direct observation of the atom. D. Scientists cannot measure the speed of electrons with sufficient accuracy to determine which model is correct. What would cause scientists to change the current model of the atom? 112 Chapter 3

42 INTERPRETING GRAPHICS Directions (10 13): For each question below, record the correct answer on a separate sheet of paper. Use the diagram below to answer questions 10 through 13. Energy of Orbitals 4f Energy n = 4 n = 3 4d 4p 3d 4s 3p 3s n = 2 2p 2s n = 1 1s 0 q w e Potassium has 19 protons. According to this diagram of energy levels, what is the energy level of the most energetic electrons in a potassium atom at its ground state? F. 1s H. 3p G. 3d I. 4s Which of these electron transitions emits the largest amount of energy? A. 2s to 3d B. 2s to 4s C. 3d to 2s D. 4s to 2s Why is the 4s level below the 3d level on this chart? F. There are ten 3d electrons but only two 4s electrons. G. The 4s electrons have lower energy than the 3d electrons. H. It is just a convention to save space when drawing the chart. I. There is a smaller transition between 4s and 3p than between 4s and 3d. The element, titanium, has two electrons in the 3d orbital. What is the atomic number of titanium? Test To develop a shortresponse or extendedresponse answer, jot down your key ideas on a piece of scratch paper first (if allowed), then expand on these ideas to build your answer. Standardized Test Prep 113